the weight of silver chloride thus obtained calculate the percentageof chlorine in the sample of sodium chloride.

[Note 1: The washed asbestos for this type of filter is prepared bydigesting in concentrated hydrochloric acid, long-fibered asbestoswhich has been cut in pieces of about 0.5 cm. in length. Afterdigestion, the asbestos is filtered off on a filter plate and washedwith hot, distilled water until free from chlorides. A small portionof the asbestos is shaken with water, forming a thin suspension, whichis bottled and kept for use.]

[Note 2: The nitric acid is added before precipitation to lessen thetendency of the silver chloride to carry down with it other substanceswhich might be precipitated from a neutral solution. A large excess ofthe acid would exert a slight solvent action upon the chloride.]

[Note 3: The solution should not be boiled after the addition of thenitric acid before the presence of an excess of silver nitrate isassured, since a slight interaction between the nitric acid and thesodium chloride is possible, by which a loss of chlorine, either assuch or as hydrochloric acid, might ensue. The presence of an excessof the precipitant can usually be recognized at the time of itsaddition, by the increased readiness with which the precipitatecoagulates and settles.]

[Note 4: The precipitate should not be exposed to strong sunlight,since under those conditions a reduction of the silver chloride ensueswhich is accompanied by a loss of chlorine. The superficial alterationwhich the chloride undergoes in diffused daylight is not sufficientto materially affect the accuracy of the determination. It should benoted, however, that a slight error does result from the effect oflight upon the silver chloride precipitate and in cases in which thegreatest obtainable accuracy is required, the procedure describedunder "Method B" should be followed, in which this slight reduction ofthe silver chloride is corrected by subsequent treatment with nitricand hydrochloric acids.]

[Note 5: The asbestos used in the Gooch filter should be of the finestquality and capable of division into minute fibrous particles. Acoarse felt is not satisfactory.]

[Note 6: The precipitate must be washed with warm water until it isabsolutely free from silver and sodium nitrates. It may be assumedthat the sodium salt is completely removed when the wash-water showsno evidence of silver. It must be borne in mind that silver chlorideis somewhat soluble in hydrochloric acid, and only a single dropshould be added. The washing should be continued until no cloudinesswhatever can be detected in 3 cc. of the washings.

Silver chloride is but slightly soluble in water. The solubilityvaries with its physical condition within small limits, and isabout 0.0018 gram per liter at 18 deg.C. for the curdy variety usuallyprecipitated. The chloride is also somewhat soluble in solutions ofmany chlorides, in solutions of silver nitrate, and in concentratednitric acid.

As a matter of economy, the filtrate, which contains whatever silvernitrate was added in excess, may be set aside. The silver can beprecipitated as chloride and later converted into silver nitrate.]

[Note 7: The use of the Gooch filter commends itself strongly when aconsiderable number of halogen determinations are to be made, sincesuccessive portions of the silver halides may be filtered on the samefilter, without the removal of the preceding portions, until thecrucible is about two thirds filled. If the felt is properly prepared,filtration and washing are rapidly accomplished on this filter, andthis, combined with the possibility of collecting several precipitateson the same filter, is a strong argument in favor of its use with anybut gelatinous precipitates.]

!Method B. With the Use of a Paper Filter!

PROCEDURE.--Weigh out two portions of sodium chloride of about0.25-0.3 gram each and proceed with the precipitation of the silverchloride as described under Method A above. When the chloride is readyfor filtration prepare two 9 cm. washed paper filters (see Appendix).Pour the liquid above the precipitates through the filters, wash twiceby decantation and transfer the precipitates to the filters, finallywashing them until free from silver solution as described. The funnelshould then be covered with a moistened filter paper by stretching itover the top and edges, to which it will adhere on drying. It shouldbe properly labeled with the student's name and desk number, and thenplaced in a drying closet, at a temperature of about 100-110 deg.C., untilcompletely dry.

The perfectly dry filter is then opened over a circular piece ofclean, smooth, glazed paper about six inches in diameter, placed upona larger piece about twelve inches in diameter. The precipitate isremoved from the filter as completely as possible by rubbing the sidesgently together, or by scraping them cautiously with a feather whichhas been cut close to the quill and is slightly stiff (Note 1). Ineither case, care must be taken not to rub off any considerablequantity of the paper, nor to lose silver chloride in the form ofdust. Cover the precipitate on the glazed paper with a watch-glass toprevent loss of fine particles and to protect it from dust from theair. Fold the filter paper carefully, roll it into a small cone, andwind loosely around !the top! a piece of small platinum wire (Note 2).Hold the filter by the wire over a small porcelain crucible (which hasbeen cleaned, ignited, cooled in a desiccator, and weighed), igniteit, and allow the ash to fall into the crucible. Place the crucibleupon a clean clay triangle, on its side, and ignite, with a lowflame well at its base, until all the carbon of the filter has beenconsumed. Allow the crucible to cool, add two drops of concentratednitric acid and one drop of concentrated hydrochloric acid, and heat!very cautiously!, to avoid spattering, until the acids have beenexpelled; then transfer the main portion of the precipitate from theglazed paper to the cooled crucible, placing the latter on the largerpiece of glazed paper and brushing the precipitate from thesmaller piece into it, sweeping off all particles belonging to thedetermination.

Moisten the precipitate with two drops of concentrated nitric acid andone drop of concentrated hydrochloric acid, and again heat with greatcaution until the acids are expelled and the precipitate is white,when the temperature is slowly raised until the silver chloride justbegins to fuse at the edges (Note 3). The crucible is then cooled in adesiccator and weighed, after which the heating (without the additionof acids) is repeated, and it is again weighed. This must be continueduntil the weight is constant within 0.0003 gram in two consecutiveweighings. Deduct the weight of the crucible, and calculate thepercentage of chlorine in the sample of sodium chloride taken foranalysis.

[Note 1: The separation of the silver chloride from the filter isessential, since the burning carbon of the paper would reduce aconsiderable quantity of the precipitate to metallic silver, and itscomplete reconversion to the chloride within the crucible, by means ofacids, would be accompanied by some difficulty. The small amount ofsilver reduced from the chloride adhering to the filter paper afterseparating the bulk of the precipitate, and igniting the paperas prescribed, can be dissolved in nitric acid, and completelyreconverted to chloride by hydrochloric acid. The subsequent additionof the two acids to the main portion of the precipitate restores thechlorine to any chloride which may have been partially reduced by thesunlight. The excess of the acids is volatilized by heating.]

[Note 2: The platinum wire is wrapped around the top of the filterduring its incineration to avoid contact with any reduced silver fromthe reduction of the precipitate. If the wire were placed nearer theapex, such contact could hardly be avoided.]

[Note 3: Silver chloride should not be heated to complete fusion,since a slight loss by volatilization is possible at hightemperatures. The temperature of fusion is not always sufficientto destroy filter shreds; hence these should not be allowed tocontaminate the precipitate.]

DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE,

FESO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O

DETERMINATION OF IRON

PROCEDURE.--Weigh out into beakers (200-250 cc.) two portions of thesample (Note 1) of about 1 gram each and dissolve these in 50 cc. ofwater, to which 1 cc. of dilute hydrochloric acid (sp. gr. 1.12) hasbeen added (Note 2). Heat the solution to boiling, and while at theboiling point add concentrated nitric acid (sp. gr. 1.42), !drop bydrop! (noting the volume used), until the brown coloration, whichappears after the addition of a part of the nitric acid, gives placeto a yellow or red (Note 3). Avoid a large excess of nitric acid, butbe sure that the action is complete. Pour this solution cautiouslyinto about 200 cc. of water, containing a slight excess of ammonia.Calculate for this purpose the amount of aqueous ammonia required toneutralize the hydrochloric and nitric acids added (see Appendix fordata), and also to precipitate the iron as ferric hydroxide from theweight of the ferrous ammonium sulphate taken for analysis, assumingit to be pure (Note 4). The volume thus calculated will be in excessof that actually required for precipitation, since the acids are inpart consumed in the oxidation process, or are volatilized. Heat thesolution to boiling, and allow the precipitated ferric hydroxide tosettle. Decant the clear liquid through a washed filter (9 cm.),keeping as much of the precipitate in the beaker as possible. Washtwice by decantation with 100 cc. of hot water. Reserve the filtrate.Dissolve the iron from the filter with hot, dilute hydrochloric acid(sp. gr. 1.12), adding it in small portions, using as little aspossible and noting the volume used. Collect the solution in thebeaker in which precipitation took place. Add 1 cc. of nitric acid(sp. gr. 1.42), boil for a few moments, and again pour into acalculated excess of ammonia.

Wash the precipitate twice by decantation, and finally transfer it tothe original filter. Wash continuously with hot water until finally3 cc. of the washings, acidified with nitric acid (Note 5), showno evidences of the presence of chlorides when tested with silvernitrate. The filtrate and washings are combined with those from thefirst precipitation and treated for the determination of sulphur, asprescribed on page 112.

[Note 1: If a selection of pure material for analysis is to be made,crystals which are cloudy are to be avoided on account of loss ofwater of crystallization; and also those which are red, indicatingthe presence of ferric iron. If, on the other hand, the value of anaverage sample of material is desired, it is preferable to grind thewhole together, mix thoroughly, and take a sample from the mixture foranalysis.]

[Note 2: When aqueous solutions of ferrous compounds are heated in theair, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs inthe absence of free acid. The H^{+} and OH^{-} ions from water areinvolved in the oxidation process and the result is, in effect, theformation of some ferric hydroxide which tends to separate. Moreover,at the boiling temperature, the ferric sulphate produced by theoxidation hydrolyzes in part with the formation of a basic ferricsulphate, which also tends to separate from solution. The addition ofthe hydrochloric acid prevents the formation of ferric hydroxide, andso far reduces the ionization of the water that the hydrolysis of theferric sulphate is also prevented, and no precipitation occurs onheating.]

[Note 3: The nitric acid, after attaining a moderate strength,oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of anintermediate nitroso-compound similar in character to that formed inthe "ring-test" for nitrates. The nitric oxide is driven out by heat,and the solution then shows by its color the presence of ferriccompounds. A drop of the oxidized solution should be tested ona watch-glass with potassium ferricyanide, to insure a completeoxidation. This oxidation of the iron is necessary, since Fe^{++} ionsare not completely precipitated by ammonia.

The ionic changes which are involved in this oxidation are perhapsmost simply expressed by the equation

3Fe^{++} + NO_{3}^{-}+ 4H^{+} --> 3Fe^{+++} + 2H_{2}O + NO,

the H^{+} ions coming from the acid in the solution, in this caseeither the nitric or the hydrochloric acid. The full equation on whichthis is based may be written thus:

assuming that only enough nitric acid is added to complete theoxidation.]

[Note 4: The ferric hydroxide precipitate tends to carry down somesulphuric acid in the form of basic ferric sulphate. This tendency islessened if the solution of the iron is added to an excess of OH^{-}ions from the ammonium hydroxide, since under these conditionsimmediate and complete precipitation of the ferric hydroxide ensues.A gradual neutralization with ammonia would result in the localformation of a neutral solution within the liquid, and subsequentdeposition of a basic sulphate as a consequence of a local deficiencyof OH^{-} ions from the NH_{4}OH and a partial hydrolysis of theferric salt. Even with this precaution the entire absence of sulphatesfrom the first iron precipitate is not assured. It is, therefore,redissolved and again thrown down by ammonia. The organic matter ofthe filter paper may occasion a partial reduction of the iron duringsolution, with consequent possibility of incomplete subsequentprecipitation with ammonia. The nitric acid is added to reoxidize thisiron.

To avoid errors arising from the solvent action of ammoniacalliquids upon glass, the iron precipitate should be filtered withoutunnecessary delay.]

[Note 5: The washings from the ferric hydroxide are acidified withnitric acid, before testing with silver nitrate, to destroy theammonia which is a solvent of silver chloride.

The use of suction to promote filtration and washing is permissible,though not prescribed. The precipitate should not be allowed to dryduring the washing.]

!Ignition of the Iron Precipitate!

Heat a platinum or porcelain crucible, cool it in a desiccator andweigh, repeating until a constant weight is obtained.

Fold the top of the filter paper over the moist precipitate of ferrichydroxide and transfer it cautiously to the crucible. Wipe the insideof the funnel with a small fragment of washed filter paper, ifnecessary, and place the paper in the crucible.

Incline the crucible on its side, on a triangle supported on aring-stand, and stand the cover on edge at the mouth of the crucible.Place a burner below the front edge of the crucible, using a low flameand protecting it from drafts of air by means of a chimney. The heatfrom the burner is thus reflected into the crucible and driesthe precipitate without danger of loss as the result of a suddengeneration of steam within the mass of ferric hydroxide. As the dryingprogresses the burner may be gradually moved toward the base of thecrucible and the flame increased until the paper of the filter beginsto char and finally to smoke, as the volatile matter is expelled. Thisis known as "smoking off" a filter, and the temperature should not beraised sufficiently high during this process to cause the paper toignite, as the air currents produced by the flame of the blazing papermay carry away particles of the precipitate.

When the paper is fully charred, move the burner to the base of thecrucible and raise the temperature to the full heat of the burner forfifteen minutes, with the crucible still inclined on its side, butwithout the cover (Note 1). Finally set the crucible upright in thetriangle, cover it, and heat at the full temperature of a blast lampor other high temperature burner. Cool and weigh in the usual manner(Note 2). Repeat the strong heating until the weight is constantwithin 0.0003 gram.

From the weight of ferric oxide (Fe_{2}O_{3}) calculate the percentageof iron (Fe) in the sample (Note 3).

[Note 1: These directions for the ignition of the precipitate must beclosely followed. A ready access of atmospheric oxygen is of specialimportance to insure the reoxidation to ferric oxide of any iron whichmay be reduced to magnetic oxide (Fe_{3}O_{4}) during the combustionof the filter. The final heating over the blast lamp is essentialfor the complete expulsion of the last traces of water from thehydroxide.]

[Note 2: Ignited ferric oxide is somewhat hygroscopic. On this accountthe weighings must be promptly completed after removal from thedesiccator. In all weighings after the first it is well to place theweights upon the balance-pan before removing the crucible from thedesiccator. It is then only necessary to move the rider to obtain theweight.]

[Note 3: The gravimetric determination of aluminium or chromium iscomparable with that of iron just described, with the additionalprecaution that the solution must be boiled until it contains but avery slight excess of ammonia, since the hydroxides of aluminium andchromium are more soluble than ferric hydroxide.

The most important properties of these hydroxides, from a quantitativestandpoint, other than those mentioned, are the following: All areprecipitable by the hydroxides of sodium and potassium, but alwaysinclose some of the precipitant, and should be reprecipitated withammonium hydroxide before ignition to oxides. Chromium and aluminiumhydroxides dissolve in an excess of the caustic alkalies and formanions, probably of the formula AlO_2^{-} and CrO_{2}^{-}. Chromiumhydroxide is reprecipitated from this solution on boiling. When firstprecipitated the hydroxides are all readily soluble in acids, butaluminium hydroxide dissolves with considerable difficulty afterstanding or boiling for some time. The precipitation of the hydroxidesis promoted by the presence of ammonium chloride, but is partiallyor entirely prevented by the presence of tartaric or citric acids,glycerine, sugars, and some other forms of soluble organic matter.The hydroxides yield on ignition an oxide suitable for weighing(Al_{2}O_{3}, Cr_{2}O_{3}, Fe_{2}O_{3}).]

DETERMINATION OF SULPHUR

PROCEDURE.--Add to the combined filtrates from the ferric hydroxideabout 0.6 gram of anhydrous sodium carbonate; cover the beaker, andthen add dilute hydrochloric acid (sp. gr. 1.12) in moderate excessand evaporate to dryness on the water bath. Add 10 cc. of concentratedhydrochloric acid (sp. gr. 1.20) to the residue, and again evaporateto dryness on the bath. Dissolve the residue in water, filter if notclear, transfer to a 700 cc. beaker, dilute to about 400 cc., andcautiously add hydrochloric acid until the solution shows a distinctlyacid reaction (Note 1). Heat the solution to boiling, and add !veryslowly! and with constant stirring, 20 cc. in excess of the calculatedamount of a hot barium chloride solution, containing about 20 gramsBaCl_{2}.2H_{2}O per liter (Notes 2 and 3). Continue the boiling forabout two minutes, allow the precipitate to settle, and decant theliquid at the end of half an hour (Note 4). Replace the beakercontaining the original filtrate by a clean beaker, wash theprecipitated sulphate by decantation with hot water, and subsequentlyupon the filter until it is freed from chlorides, testing the washingsas described in the determination of iron. The filter is thentransferred to a platinum or porcelain crucible and ignited, asdescribed above, until the weight is constant (Note 5). From theweight of barium sulphate (BaSO_{4}) obtained, calculate thepercentage of sulphur (S) in the sample.

[Note 1: Barium sulphate is slightly soluble in hydrochloric acid,even dilute, probably as a result of the reduction in the degree ofdissociation of sulphuric acid in the presence of the H^{+} ions ofthe hydrochloric acid, and possibly because of the formation of acomplex anion made up of barium and chlorine; hence only the smallestexcess should be added over the amount required to acidify thesolution.]

[Note 2: The ionic changes involved in the precipitation of bariumsulphate are very simple:

Ba^{++} + SO_{4}^{--} --> [BaSO_{4}]

This case affords one of the best illustrations of the effect of anexcess of a precipitant in decreasing the solubility of a precipitate.If the conditions are considered which exist at the moment when justenough of the Ba^{++} ions have been added to correspond to theSO_{4}^{--} ions in the solution, it will be seen that nearly all ofthe barium sulphate has been precipitated, and that the small amountwhich then remains in the solution which is in contact with theprecipitate must represent a saturated solution for the existingtemperature, and that this solution is comparable with a solution ofsugar to which more sugar has been added than will dissolve. Itshould be borne in mind that the quantity of barium sulphate inthis !saturated solution is a constant quantity! for the existingconditions. The dissolved barium sulphate, like any electrolyte, isdissociated, and the equilibrium conditions may be expressed thus:

(!Conc'n Ba^{++} x Conc'n SO_{4}^{--})/(Conc'n BaSO_{4}) = Const.!,

and since !Conc'n BaSO_{4}! for the saturated solution has a constantvalue (which is very small), it may be eliminated, when the expressionbecomes !Conc'n Ba^{++} x Conc'n SO_{4}^{--} = Const.!, which isthe "solubility product" of BaSO_{4}. If, now, an excess of theprecipitant, a soluble barium salt, is added in the form of arelatively concentrated solution (the slight change of volume of a fewcubic centimeters may be disregarded for the present discussion)the concentration of the Ba^{++} ions is much increased, and as aconsequence the !Conc'n SO_{4}! must decrease in proportion if thevalue of the expression is to remain constant, which is a requisitecondition if the law of mass action upon which our argument dependsholds true. In other words, SO_{4}^{--} ions must combine with someof the added Ba^{++} ions to form [BaSO_{4}]; but it will be recalledthat the solution is already saturated with BaSO_{4}, and this freshlyformed quantity must, therefore, separate and add itself to theprecipitate. This is exactly what is desired in order to insuremore complete precipitation and greater accuracy, and leads to theconclusion that the larger the excess of the precipitant added themore successful the analysis; but a practical limit is placed uponthe quantity of the precipitant which may be properly added by otherconditions, as stated in the following note.]

[Note 3: Barium sulphate, in a larger measure than most compounds,tends to carry down other substances which are present in the solutionfrom which it separates, even when these other substances arerelatively soluble, and including the barium chloride used as theprecipitant. This is also notably true in the case of nitrates andchlorates of the alkalies, and of ferric compounds; and, since in thisanalysis ammonium nitrate has resulted from the neutralization of theexcess of the nitric acid added to oxidize the iron, it is essentialthat this should be destroyed by repeated evaporation with arelatively large quantity of hydrochloric acid. During evaporation amutual decomposition of the two acids takes place, and the nitric acidis finally decomposed and expelled by the excess of hydrochloric acid.

Iron is usually found in the precipitate of barium sulphate whenthrown down from hot solutions in the presence of ferric salts. This,according to Kuster and Thiel (!Zeit. anorg. Chem.!, 22, 424), is dueto the formation of a complex ion (Fe(SO_{4})_{2}) which precipitateswith the Ba^{++} ion, while Richards (!Zeit. anorg. Chem.!, 23, 383)ascribes it to hydrolytic action, which causes the formation of abasic ferric complex which is occluded in the barium precipitate.Whatever the character of the compound may be, it has been shown thatit loses sulphuric anhydride upon ignition, causing low results, eventhough the precipitate contains iron.

The contamination of the barium sulphate by iron is much less in thepresence of ferrous than ferric salts. If, therefore, the sulphuralone were to be determined in the ferrous ammonium sulphate, theprecipitation by barium might be made directly from an aqueoussolution of the salt, which had been made slightly acid withhydrochloric acid.]

[Note 4: The precipitation of the barium sulphate is probably completeat the end of a half-hour, and the solution may safely be filtered atthe expiration of that time if it is desired to hasten the analysis.

As already noted, many precipitates of the general character of thissulphate tend to grow more coarsely granular if digested for some timewith the liquid from which they have separated. It is therefore wellto allow the precipitate to stand in a warm place for several hours,if practicable, to promote ease of filtration. The filtrate andwashings should always be carefully examined for minute quantities ofthe sulphate which may pass through the pores of the filter. This isbest accomplished by imparting to the filtrate a gentle rotary motion,when the sulphate, if present, will collect at the center of thebottom of the beaker.]

[Note 5: A reduction of barium sulphate to the sulphide may veryreadily be caused by the reducing action of the burning carbon of thefilter, and much care should be taken to prevent any considerablereduction from this cause. Subsequent ignition, with ready accessof air, reconverts the sulphide to sulphate unless a considerablereduction has occurred. In the latter case it is expedient to add oneor two drops of sulphuric acid and to heat cautiously until the excessof acid is expelled.]

[Note 6: Barium sulphate requires about 400,000 parts of water forits solution. It is not decomposed at a red heat but suffers loss,probably of sulphur trioxide, at a temperature above 900 deg.C.]

DETERMINATION OF SULPHUR IN BARIUM SULPHATE

PROCEDURE.--Weigh out, into platinum crucibles, two portions of about0.5 gram of the sulphate. Mix each in the crucible with five to sixtimes its weight of anhydrous sodium carbonate. This can best be doneby placing the crucible on a piece of glazed paper and stirring themixture with a clean, dry stirring-rod, which may finally be wiped offwith a small fragment of filter paper, the latter being placed in thecrucible. Cover the crucible and heat until a quiet, liquid fusionensues. Remove the burner, and tip the crucible until the fused massflows nearly to its mouth. Hold it in that position until the mass hassolidified. When cold, the material may usually be detached in a lumpby tapping the crucible or gently pressing it near its upper edge. Ifit still adheres, a cubic centimeter or so of water may be placed inthe cold crucible and cautiously brought to boiling, when the cakewill become loosened and may be removed and placed in about 250 cc.of hot, distilled water to dissolve. Clean the crucible completely,rubbing the sides with a rubber-covered stirring-rod, if need be.

When the fused mass has completely disintegrated and nothing furtherwill dissolve, decant the solution from the residue of bariumcarbonate (Note 1). Pour over the residue 20 cc. of a solution ofsodium carbonate and 10 cc. of water and heat to gentle boiling forabout three minutes (Note 2). Filter off the carbonate and wash itwith hot water, testing the slightly acidified washings for sulphateand preserving any precipitates which appear in these tests. Acidifythe filtrate with hydrochloric acid until just acid, bring to boiling,and slowly add hot barium chloride solution, as in the precedingdetermination. Add also any tests from the washings in whichprecipitates have appeared. Filter, wash, ignite, and weigh.

From the weight of barium sulphate, calculate the percentage ofsulphur (S) in the sample.

[Note 1: This alkaline fusion is much employed to disintegratesubstances ordinarily insoluble in acids into two components, oneof which is water soluble and the other acid soluble. The reactioninvolved is:

BaSO_{4} + Na_{2}CO_{3}, --> BaCO_{3}, + Na_{2}SO_{4}.

As the sodium sulphate is soluble in water, and the barium carbonateinsoluble, a separation between them is possible and the sulphur canbe determined in the water-soluble portion.

It should be noted that this method can be applied to the purificationof a precipitate of barium sulphate if contaminated by most of thesubstances mentioned in Note 3 on page 114. The impurities pass intothe water solution together with the sodium sulphate, but, beingpresent in such minute amounts, do not again precipitate with thebarium sulphate.]

[Note 2: The barium carbonate is boiled with sodium carbonate solutionbefore filtration because the reaction above is reversible; and it isonly by keeping the sodium carbonate present in excess until nearlyall of the sodium sulphate solution has been removed by filtrationthat the reversion of some of the barium carbonate to barium sulphateis prevented. This is an application of the principle of mass action,in which the concentration of the reagent (the carbonate ion) iskept as high as practicable and that of the sulphate ion as low aspossible, in order to force the reaction in the desired direction (seeAppendix).]

DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE

The mineral apatite is composed of calcium phosphate, associated withcalcium chloride, or fluoride. Specimens are easily obtainable whichare nearly pure and leave on treatment with acid only a slightsiliceous residue.

For the purpose of gravimetric determination, phosphoric acid isusually precipitated from ammoniacal solutions in the form ofmagnesium ammonium phosphate which, on ignition, is converted intomagnesium pyrophosphate. Since the calcium phosphate of the apatiteis also insoluble in ammoniacal solutions, this procedure cannot beapplied directly. The separation of the phosphoric acid from thecalcium must first be accomplished by precipitation in the form ofammonium phosphomolybdate in nitric acid solution, using ammoniummolybdate as the precipitant. The "yellow precipitate," as it is oftencalled, is not always of a definite composition, and therefore notsuitable for direct weighing, but may be dissolved in ammonia, and thephosphoric acid thrown out as magnesium ammonium phosphate from thesolution.

Of the substances likely to occur in apatite, silicic acid aloneinterferes with the precipitation of the phosphoric acid in nitricacid solution.

PRECIPITATION OF AMMONIUM PHOSPHOMOLYBDATE

PROCEDURE.--Grind the mineral in an agate mortar until no grit isperceptible. Transfer the substance to a weighing-tube, and weigh outtwo portions, not exceeding 0.20 gram each (Note 1) into two beakersof about 200 cc. capacity. Pour over them 20 cc. of dilute nitric acid(sp. gr. 1.2) and warm gently until solvent action has apparentlyceased. Evaporate the solution cautiously to dryness, heat the residuefor about an hour at 100-110 deg.C., and treat it again with nitric acidas described above; separate the residue of silica by filtration ona small filter (7 cm.) and wash with warm water, using as little aspossible (Note 2). Receive the filtrate in a beaker (200-500 cc.).Test the washings with ammonia for calcium phosphate, but add all suchtests in which a precipitate appears to the original nitrate (Note 3).The filtrate and washings must be kept as small as possible and shouldnot exceed 100 cc. in volume. Add aqueous ammonia (sp. gr. 0.96) untilthe precipitate of calcium phosphate first produced just fails toredissolve, and then add a few drops of nitric acid until this isagain brought into solution (Note 4). Warm the solution until itcannot be comfortably held in the hand (about 60 deg.C.) and, afterremoval of the burner, add 75 cc. of ammonium molybdate solution whichhas been !gently! warmed, but which must be perfectly clear. Allowthe mixture to stand at a temperature of about 50 or 60 deg.C. for twelvehours (Notes 5 and 6). Filter off the yellow precipitate on a 9 cm.filter, and wash by decantation with a solution of ammonium nitratemade acid with nitric acid.[1] Allow the precipitate to remain in thebeaker as far as possible. Test the washings for calcium with ammoniaand ammonium oxalate (Note 3).

Add 10 cc. of molybdate solution to the nitrate, and leave it fora few hours. It should then be carefully examined for a !yellow!precipitate; a white precipitate may be neglected.

[Note 1: Magnesium ammonium phosphate, as noted below, is slightlysoluble under the conditions of operation. Consequently theunavoidable errors of analysis are greater in this determination thanin those which have preceded it, and some divergence may be expectedin duplicate analyses. It is obvious that the larger the amount ofsubstance taken for analysis the less will be the relative loss orgain due to unavoidable experimental errors; but, in this instance, acheck is placed upon the amount of material which may be taken both bythe bulk of the resulting precipitate of ammonium phosphomolybdateand by the excessive amount of ammonium molybdate required to effectcomplete separation of the phosphoric acid, since a liberal excessabove the theoretical quantity is demanded. Molybdic acid is one ofthe more expensive reagents.]

[Note 2: Soluble silicic acid would, if present, partially separatewith the phosphomolybdate, although not in combination withmolybdenum. Its previous removal by dehydration is thereforenecessary.]

[Note 3: When washing the siliceous residue the filtrate may be testedfor calcium by adding ammonia, since that reagent neutralizes theacid which holds the calcium phosphate in solution and causesprecipitation; but after the removal of the phosphoric acid incombination with the molybdenum, the addition of an oxalate isrequired to show the presence of calcium.]

[Note 4: An excess of nitric acid exerts a slight solventaction, while ammonium nitrate lessens the solubility; hence theneutralization of the former by ammonia.]

[Note 5: The precipitation of the phosphomolybdate takes place morepromptly in warm than in cold solutions, but the temperature shouldnot exceed 60 deg.C. during precipitation; a higher temperature tends toseparate molybdic acid from the solution. This acid is nearly white,and its deposition in the filtrate on long standing should not bemistaken for a second precipitation of the yellow precipitate. Theaddition of 75 cc. of ammonium molybdate solution insures the presenceof a liberal excess of the reagent, but the filtrate should be testedas in all quantitative procedures.

The precipitation is probably complete in many cases in less thantwelve hours; but it is better, when practicable, to allow thesolution to stand for this length of time. Vigorous shaking orstirring promotes the separation of the precipitate.]

[Note 6: The composition of the "yellow precipitate" undoubtedlyvaries slightly with varying conditions at the time of its formation.Its composition may probably fairly be represented by the formula,(NH_{4})_{3}PO_{4}.12MoO_{3}.H_{2}O, when precipitated under theconditions prescribed in the procedure. Whatever other variations mayoccur in its composition, the ratio of 12 MoO_{3}:1 P seems tohold, and this fact is utilized in volumetric processes for thedetermination of phosphorus, in which the molybdenum is reduced toa lower oxide and reoxidized by a standard solution of potassiumpermanganate. In principle, the procedure is comparable with thatdescribed for the determination of iron by permanganate.]

PRECIPITATION OF MAGNESIUM AMMONIUM PHOSPHATE

PROCEDURE.--Dissolve the precipitate of phosphomolybdate upon thefilter by pouring through it dilute aqueous ammonia (one volume ofdilute ammonia (sp. gr. 0.96) and three volumes of water, whichshould be carefully measured), and receive the solution in the beakercontaining the bulk of the precipitate. The total volume of nitrateand washings should not much exceed 100 cc. Acidify the solution withdilute hydrochloric acid, and heat it nearly to boiling. Calculate thevolume of magnesium ammonium chloride solution ("magnesia mixture")required to precipitate the phosphoric acid, assuming 40 per centP_{2}O_{5} in the apatite. Measure out about 5 cc. in excess of thisamount, and pour it into the acid solution. Then add slowly diluteammonium hydroxide (1 volume of strong ammonia (sp. gr. 0.90) and 9volumes of water), stirring constantly until a precipitate forms. Thenadd a volume of filtered, concentrated ammonia (sp. gr. 0.90) equal toone third of the volume of liquid in the beaker (Note 1). Allow thewhole to cool. The precipitated magnesium ammonium phosphate shouldthen be definitely crystalline in appearance (Note 2). (If it isdesired to hasten the precipitation, the solution may be cooled, firstin cold and then in ice-water, and stirred !constantly! for half anhour, when precipitation will usually be complete.)

Decant the clear liquid through a filter, and transfer the precipitateto the filter, using as wash-water a mixture of one volume ofconcentrated ammonia and three volumes of water. It is not necessaryto clean the beaker completely or to wash the precipitate thoroughlyat this point, as it is necessary to purify it by reprecipitation.

[Note 1: Magnesium ammonium phosphate is not a wholly insolublesubstance, even under the most favorable analytical conditions. Itis least soluble in a liquid containing one fourth of its volume ofconcentrated aqueous ammonia (sp. gr. 0.90) and this proportion shouldbe carefully maintained as prescribed in the procedure. On account ofthis slight solubility the volume of solutions should be kept as smallas possible and the amount of wash-water limited to that absolutelyrequired.

A large excess of the magnesium solution tends both to throw outmagnesium hydroxide (shown by a persistently flocculent precipitate)and to cause the phosphate to carry down molybdic acid. The tendencyof the magnesium precipitate to carry down molybdic acid is alsoincreased if the solution is too concentrated. The volume should notbe less than 90 cc., nor more than 125 cc., at the time of the firstprecipitation with the magnesia mixture.]

[Note 2: The magnesium ammonium phosphate should be perfectlycrystalline, and will be so if the directions are followed. The slowaddition of the reagent is essential, and the stirring not less so.Stirring promotes the separation of the precipitate and the formationof larger crystals, and may therefore be substituted for digestion inthe cold. The stirring-rod must not be allowed to scratch the glass,as the crystals adhere to such scratches and are removed withdifficulty.]

REPRECIPITATION AND IGNITION OF MAGNESIUM AMMONIUM PHOSPHATE

A single precipitation of the magnesium compound in the presence ofmolybdenum compounds rarely yields a pure product. The molybdenum canbe removed by solution of the precipitate in acid and precipitationof the molybdenum by sulphureted hydrogen, after which the magnesiumprecipitate may be again thrown down. It is usually more satisfactoryto dissolve the magnesium precipitate and reprecipitate the phosphateas magnesium ammonium phosphate as described below.

PROCEDURE.--Dissolve the precipitate from the filter in a littledilute hydrochloric acid (sp. gr. 1.12), allowing the acid solution torun into the beaker in which the original precipitation was made (Note1). Wash the filter with water until the wash-water shows no test forchlorides, but avoid an unnecessary amount of wash-water. Add tothe solution 2 cc. (not more) of magnesia mixture, and then diluteammonium hydroxide solution (sp. gr. 0.96), drop by drop, withconstant stirring, until the liquid smells distinctly of ammonia. Stirfor a few moments and then add a volume of strong ammonia (sp. gr.0.90), equal to one third of the volume of the solution. Allow thesolution to stand for some hours, and then filter off the magnesiumammonium phosphate, which should be distinctly crystalline incharacter. Wash the precipitate with dilute ammonia water, asprescribed above, until, finally, 3 cc. of the washings, afteracidifying with nitric acid, show no evidence of chlorides. Test bothfiltrates for complete precipitation by adding a few cubic centimetersof magnesia mixture and allowing them to stand for some time.

Transfer the moist precipitate to a weighed porcelain or platinumcrucible and ignite, using great care to raise the temperature slowlywhile drying the filter in the crucible, and to insure the readyaccess of oxygen during the combustion of the filter paper, thusguarding against a possible reduction of the phosphate, which wouldresult in disastrous consequences both to the crucible, if ofplatinum, and the analysis. Do not raise the temperature abovemoderate redness until the precipitate is white. (Keep this precautionwell in mind.) Ignite finally at the highest temperature of theTirrill burner, and repeat the heating until the weight is constant.If the ignited precipitate is persistently discolored by particles ofunburned carbon, moisten the mass with a drop or two of concentratednitric acid and heat cautiously, finally igniting strongly. Theacid will dissolve magnesium pyrophosphate from the surface of theparticles of carbon, which will then burn away. Nitric acid also aidsas an oxidizing agent in supplying oxygen for the combustion of thecarbon.

From the weight of magnesium pyrophosphate (Mg_{2}P_{2}O_{7})obtained, calculate the phosphoric anhydride (P_{2}O_{5}) in thesample of apatite.

[Note 1: The ionic change involved in the precipitation of themagnesium compound is

PO_{4}^{---} + NH_{4}^{+} + Mg^{++} --> [MgNH_{4}PO_{4}].

The magnesium ammonium phosphate is readily dissolved by acids, eventhose which are no stronger than acetic acid. This is accounted forby the fact that two of the ions into which phosphoric acid maydissociate, the HPO_{4}^{--} or H_{2}PO_{4}^{-} ions, exhibit thecharacteristics of very weak acids, in that they show almost notendency to dissociate further into H^{+} and PO_{4}^{--} ions.Consequently the ionic changes which occur when the magnesium ammoniumphosphate is brought into contact with an acid may be typified by thereaction:

that is, the PO_{4}^{--} ions and the H^{+} ions lose their identityin the formation of the new ion, HPO_{4}^{--}, and this continuesuntil the magnesium ammonium phosphate is entirely dissolved.]

[Note 2: During ignition the magnesium ammonium phosphate losesammonia and water and is converted into magnesium pyrophosphate:

2MgNH_{4}PO_{4} --> Mg_{2}P_{2}O_{7} + 2NH_{3} + H_{2}O.

The precautions mentioned on pages 111 and 123 must be observed withgreat care during the ignition of this precipitate. The danger herelies in a possible reduction of the phosphate by the carbon of thefilter paper, or by the ammonia evolved, which may act as a reducingagent. The phosphorus then attacks and injures a platinum crucible,and the determination is valueless.]

ANALYSIS OF LIMESTONE

Limestones vary widely in composition from a nearly pure marblethrough the dolomitic limestones, containing varying amounts ofmagnesium, to the impure varieties, which contain also ferrous andmanganous carbonates and siliceous compounds in variable proportions.Many other minerals may be inclosed in limestones in small quantities,and an exact qualitative analysis will often show the presence ofsulphides or sulphates, phosphates, and titanates, and the alkali oreven the heavy metals. No attempt is made in the following proceduresto provide a complete quantitative scheme which would take intoaccount all of these constituents. Such a scheme for a completeanalysis of a limestone may be found in Bulletin No. 700 of the UnitedStates Geological Survey. It is assumed that, for these practicedeterminations, a limestone is selected which contains only the morecommon constituents first enumerated above.

DETERMINATION OF MOISTURE

The determination of the amount of moisture in minerals or ores isoften of great importance. Ores which have been exposed to the weatherduring shipment may have absorbed enough moisture to appreciablyaffect the results of analysis. Since it is essential that the sellerand buyer should make their analyses upon comparable material, it iscustomary for each analyst to determine the moisture in the sampleexamined, and then to calculate the percentages of the variousconstituents with reference to a sample dried in the air, or at atemperature a little above 100 deg.C., which, unless the ore has undergonechemical change because of the wetting, should be the same before andafter shipment.

PROCEDURE.--Spread 25 grams of the powdered sample on a weighedwatch-glass; weigh to the nearest 10 milligrams only and heat at105 deg.C.; weigh at intervals of an hour, after cooling in a desiccator,until the loss of weight after an hour's heating does not exceed10 milligrams. It should be noted that a variation in weight of 10milligrams in a total weight of 25 grams is no greater relatively thana variation of 0.1 milligram when the sample taken weighs 0.25 gram

DETERMINATION OF THE INSOLUBLE MATTER AND SILICA

PROCEDURE.--Weigh out two portions of the original powdered sample(not the dried sample), of about 5 grams each, into 250 cc.casseroles, and cover each with a watch-glass (Note 1). Pour over thepowder 25 cc. of water, and then add 50 cc. of dilute hydrochloricacid (sp. gr. 1.12) in small portions, warming gently, until nothingfurther appears to dissolve (Note 2). Evaporate to dryness on thewater bath. Pour over the residue a mixture of 5 cc. of water and5 cc. of concentrated hydrochloric acid (sp. gr. 1.2) and againevaporate to dryness, and finally heat for at least an hour ata temperature of 110 deg.C. Pour over this residue 50 cc. of dilutehydrochloric acid (one volume acid (sp. gr. 1.12) to five volumeswater), and boil for about five minutes; then filter and wash twicewith the dilute hydrochloric acid, and then with hot water untilfree from chlorides. Transfer the filter and contents to a porcelaincrucible, dry carefully over a low flame, and ignite to constantweight. The residue represents the insoluble matter and the silicafrom any soluble silicates (Note 3).

Calculate the combined percentage of these in the limestone.

[Note 1: The relatively large weight (5 grams) taken for analysisinsures greater accuracy in the determination of the ingredients whichare present in small proportions, and is also more likely to be arepresentative sample of the material analyzed.]

[Note 2: It is plain that the amount of the insoluble residue and alsoits character will often depend upon the strength of acid used forsolution of the limestone. It cannot, therefore, be regarded asrepresenting any well-defined constituent, and its determination isessentially empirical.]

[Note 3: It is probable that some of the silicates present are whollyor partly decomposed by the acid, and the soluble silicic acid mustbe converted by evaporation to dryness, and heating, into white,insoluble silica. This change is not complete after one evaporation.The heating at a temperature somewhat higher than that of the waterbath for a short time tends to leave the silica in the form of apowder, which promotes subsequent filtration. The siliceous residueis washed first with dilute acid to prevent hydrolytic changes, whichwould result in the formation of appreciable quantities of insolublebasic iron or aluminium salts on the filter when washing with hotwater.

If it is desired to determine the percentage of silica separately, theignited residue should be mixed in a platinum crucible with about sixtimes its weight of anhydrous sodium carbonate, and the proceduregiven on page 151 should be followed. The filtrate from the silica isthen added to the main filtrate from the insoluble residue.]

DETERMINATION OF FERRIC OXIDE AND ALUMINIUM OXIDE (WITH MANGANESE)

PROCEDURE.--To the filtrate from the insoluble residue add ammoniumhydroxide until the solution just smells distinctly of ammonia, but donot add an excess. Then add 5 cc. of saturated bromine water (Note 1),and boil for five minutes. If the smell of ammonia has disappeared,again add ammonium hydroxide in slight excess, and 3 cc. of brominewater, and heat again for a few minutes. Finally add 10 cc. ofammonium chloride solution and keep the solution warm until it barelysmells of ammonia; then filter promptly (Note 2). Wash the filtertwice with hot water, then (after replacing the receiving beaker) pourthrough it 25 cc. of hot, dilute hydrochloric acid (one volume diluteHCl [sp. gr. 1.12] to five volumes water). A brown residue insolublein the acid may be allowed to remain on the filter. Wash the filterfive times with hot water, add to the filtrate ammonium hydroxideand bromine water as described above, and repeat the precipitation.Collect the precipitate on the filter already used, wash it free fromchlorides with hot water, and ignite and weigh as described for ferrichydroxide on page 110. The residue after ignition consists of ferricoxide, alumina, and mangano-manganic oxide (Mn_{3}O_{4}), if manganeseis present. These are commonly determined together (Note 3).

Calculate the percentage of the combined oxides in the limestone.

[Note 1: The addition of bromine water to the ammoniacal solutionsserves to oxidize any ferrous hydroxide to ferric hydroxide and toprecipitate manganese as MnO(OH)_{2}. The solution must contain notmore than a bare excess of hydroxyl ions (ammonium hydroxide) when itis filtered, on account of the tendency of the aluminium hydroxide toredissolve.

The solution should not be strongly ammoniacal when the bromine isadded, as strong ammonia reacts with the bromine, with the evolutionof nitrogen.]

[Note 2: The precipitate produced by ammonium hydroxide and bromineshould be filtered off promptly, since the alkaline solution absorbscarbon dioxide from the air, with consequent partial precipitationof the calcium as carbonate. This is possible even under the mostfavorable conditions, and for this reason the iron precipitate isredissolved and again precipitated to free it from calcium. When theprecipitate is small, this reprecipitation may be omitted.]

[Note 3: In the absence of significant amounts of manganese the ironand aluminium may be separately determined by fusion of the mixedignited precipitate, after weighing, with about ten times its weightof acid potassium sulphate, solution of the cold fused mass in water,and volumetric determination of the iron, as described on page 66.The aluminium is then determined by difference, after subtracting theweight of ferric oxide corresponding to the amount of iron found.

If a separate determination of the iron, aluminium, and manganeseis desired, the mixed precipitate may be dissolved in acid beforeignition, and the separation effected by special methods (see, forexample, Fay, !Quantitative Analyses!, First Edition, pp. 15-19 and23-27).]

DETERMINATION OF CALCIUM

PROCEDURE.--To the combined filtrates from the double precipitation ofthe hydroxides just described, add 5 cc. of dilute ammonium hydroxide(sp. gr. 0.96), and transfer the liquid to a 500 cc. graduated flask,washing out the beaker carefully. Cool to laboratory temperature, andfill the flask with distilled water until the lowest point of themeniscus is exactly level with the mark on the neck of the flask.Carefully remove any drops of water which are on the inside of theneck of the flask above the graduation by means of a strip of filterpaper, make the solution uniform by pouring it out into a dry beakerand back into the flask several times. Measure off one fifth of thissolution as follows (Note 1): Pour into a 100 cc. graduated flaskabout 10 cc. of the solution, shake the liquid thoroughly over theinner surface of the small flask, and pour it out. Repeat the sameoperation. Fill the 100 cc. flask until the lowest point of themeniscus is exactly level with the mark on its neck, remove any dropsof solution from the upper part of the neck with filter paper, andpour the solution into a beaker (400-500 cc.). Wash out the flask withsmall quantities of water until it is clean, adding these to the 100cc. of solution. When the duplicate portion of 100 cc. is measured outfrom the solution, remember that the flask must be rinsed out twicewith that solution, as prescribed above, before the measurement ismade. (A 100 cc. pipette may be used to measure out the aliquotportions, if preferred.)

Dilute each of the measured portions to 250 cc. with distilled water,heat the whole to boiling, and add ammonium oxalate solution slowlyin moderate excess, stirring well. Boil for two minutes; allow theprecipitated calcium oxalate to settle for a half-hour, and decantthrough a filter. Test the filtrate for complete precipitation byadding a few cubic centimeters of the precipitant, allowing it tostand for fifteen minutes. If no precipitate forms, make the solutionslightly acid with hydrochloric acid (Note 2); see that it is properlylabeled and reserve it to be combined with the filtrate from thesecond calcium oxalate precipitation (Notes 3 and 4).

Redissolve the calcium oxalate in the beaker with warm hydrochloricacid, pouring the acid through the filter. Wash the filter five timeswith water, and finally pour through it aqueous ammonia. Dilute thesolution to 250 cc., bring to boiling, and add 1 cc. ammonium oxalatesolution (Note 5) and ammonia in slight excess; boil for two minutes,and set aside for a half-hour. Filter off the calcium oxalate upon thefilter first used, and wash free from chlorides. The filtrate shouldbe made barely acid with hydrochloric acid and combined with thefiltrate from the first precipitation. Begin at once the evaporationof the solutions for the determination of magnesium as describedbelow.

The precipitate of calcium oxalate may be converted into calcium oxideby ignition without previous drying. After burning the filter, it maybe ignited for three quarters of an hour in a platinum crucible atthe highest heat of the Bunsen or Tirrill burner, and finally for tenminutes at the blast lamp (Note 6). Repeat the heating over the blastlamp until the weight is constant. As the calcium oxide absorbsmoisture from the air, it must (after cooling) be weighed as rapidlyas possible.

The precipitate may, if preferred, be placed in a weighted porcelaincrucible. After burning off the filter and heating for ten minutes thecalcium precipitate may be converted into calcium sulphate by placing2 cc. of dilute sulphuric acid in the crucible (cold), heating thecovered crucible very cautiously over a low flame to drive off theexcess of acid, and finally at redness to constant weight (Note 7).

From the weight of the oxide or sulphate, calculate the percentage ofthe calcium (Ca) in the limestone, remembering that only one fifth ofthe total solution is used for this determination.

[Note 1: If the calcium were precipitated from the entire solution,the quantity of the precipitate would be greater than could beproperly treated. The solution is, therefore, diluted to a definitevolume (500 cc.), and exactly one fifth (100 cc.) is measured off in agraduated flask or by means of a pipette.]

[Note 2: The filtrate from the calcium oxalate should be made slightlyacid immediately after filtration, in order to avoid the solventaction of the alkaline liquid upon the glass.]

[Note 3: The accurate quantitative separation of calcium and magnesiumas oxalates requires considerable care. The calcium precipitateusually carries down with it some magnesium, and this can bestbe removed by redissolving the precipitate after filtration, andreprecipitation in the presence of only the small amount of magnesiumwhich was included in the first precipitate. When, however, theproportion of magnesium is not very large, the second precipitation ofthe calcium can usually be avoided by precipitating it from a ratherdilute solution (800 cc. or so) and in the presence of a considerableexcess of the precipitant, that is, rather more than enough to convertboth the magnesium and calcium into oxalates.]

[Note 4: The ionic changes involved in the precipitation of calciumas oxalate are exceedingly simple, and the principles discussed inconnection with the barium sulphate precipitation on page 113 alsoapply here. The reaction is

C_{2}O_{4}^{--} + Ca^{++} --> [CaC_{2}O_{4}].

Calcium oxalate is nearly insoluble in water, and only very slightlysoluble in acetic acid, but is readily dissolved by the strong mineralacids. This behavior with acids is explained by the fact that oxalicacid is a stronger acid than acetic acid; when, therefore, the oxalateis brought into contact with the latter there is almost no tendency todiminish the concentration of C_{2}O_{4}^{--} ions by the formation ofan acid less dissociated than the acetic acid itself, and practicallyno solvent action ensues. When a strong mineral acid is present,however, the ionization of the oxalic acid is much reduced by the highconcentration of the H^{+} ions from the strong acid, the formationof the undissociated acid lessens the concentration of theC_{2}O_{4}^{--} ions in solution, more of the oxalate passes intosolution to re-establish equilibrium, and this process repeats itselfuntil all is dissolved.

The oxalate is immediately reprecipitated from such a solution on theaddition of OH^{-} ions, which, by uniting with the H^{+} ions of theacids (both the mineral acid and the oxalic acid) to form water, leavethe Ca^{++} and C_{2}O_{4}^{--} ions in the solution to recombine toform [CaC_{2}O_{4}], which is precipitated in the absence of theH^{+} ions. It is well at this point to add a small excess ofC_{2}O_{4}^{--} ions in the form of ammonium oxalate to decrease thesolubility of the precipitate.

[Note 5: The small quantity of ammonium oxalate solution is addedbefore the second precipitation of the calcium oxalate to insurethe presence of a slight excess of the reagent, which promotes theseparation of the calcium compound.]

For small weights of the oxalate (0.6 gram or less) this reaction maybe brought about in a platinum crucible at the highest temperature ofa Tirrill burner, but it is well to ignite larger quantities than thisover the blast lamp until the weight is constant.]

[Note 7: The heat required to burn the filter, and that subsequentlyapplied as described, will convert most of the calcium oxalate tocalcium carbonate, which is changed to sulphate by the sulphuric acid.The reactions involved are

If a porcelain crucible is employed for ignition, this conversion tosulphate is to be preferred, as a complete conversion to oxide isdifficult to accomplish.]

[Note 8: The determination of the calcium may be completedvolumetrically by washing the calcium oxalate precipitate fromthe filter into dilute sulphuric acid, warming, and titratingthe liberated oxalic acid with a standard solution of potassiumpermanganate as described on page 72. When a considerable number ofanalyses are to be made, this procedure will save much of the timeotherwise required for ignition and weighing.]

DETERMINATION OF MAGNESIUM

PROCEDURE.--Evaporate the acidified filtrates from the calciumprecipitates until the salts begin to crystallize, but do !not!evaporate to dryness (Note 1). Dilute the solution cautiously untilthe salts are brought into solution, adding a little acid if thesolution has evaporated to very small volume. The solution should becarefully examined at this point and must be filtered if a precipitatehas appeared. Heat the clear solution to boiling; remove the burnerand add 25 cc. of a solution of disodium phosphate. Then add slowlydilute ammonia (1 volume strong ammonia (sp. gr. 0.90) and 9 volumeswater) as long as a precipitate continues to form. Finally, add avolume of concentrated ammonia (sp. gr. 0.90) equal to one third ofthe volume of the solution, and allow the whole to stand for abouttwelve hours.

Decant the solution through a filter, wash it with dilute ammoniawater, proceeding as prescribed for the determination of phosphoricanhydride on page 122, including; the reprecipitation (Note 2),except that 3 cc. of disodium phosphate solution are added before thereprecipitation of the magnesium ammonium phosphate instead ofthe magnesia mixture there prescribed. From the weight of thepyrophosphate, calculate the percentage of magnesium oxide (MgO) inthe sample of limestone. Remember that the pyrophosphate finallyobtained is from one fifth of the original sample.

[Note 1: The precipitation of the magnesium should be made in as smallvolume as possible, and the ratio of ammonia to the total volume ofsolution should be carefully provided for, on account of the relativesolubility of the magnesium ammonium phosphate. This matter hasbeen fully discussed in connection with the phosphoric anhydridedetermination.]

[Note 2: The first magnesium ammonium phosphate precipitate is rarelywholly crystalline, as it should be, and is not always of the propercomposition when precipitated in the presence of such large amounts ofammonium salts. The difficulty can best be remedied by filtering theprecipitate and (without washing it) redissolving in a small quantityof hydrochloric acid, from which it may be again thrown down byammonia after adding a little disodium phosphate solution. If theflocculent character was occasioned by the presence of magnesiumhydroxide, the second precipitation, in a smaller volume containingfewer salts, will often result more favorably.

The removal of iron or alumina from a contaminated precipitate isa matter involving a long procedure, and a redetermination of themagnesium from a new sample, with additional precautions, is usuallyto be preferred.]

DETERMINATION OF CARBON DIOXIDE

!Absorption Apparatus!

[Illustration: Fig. 3]

The apparatus required for the determination of the carbon dioxideshould be arranged as shown in the cut (Fig. 3). The flask (A) isan ordinary wash bottle, which should be nearly filled with dilutehydrochloric acid (100 cc. acid (sp. gr. 1.12) and 200 cc. of water).The flask is connected by rubber tubing (a) with the glass tube (b)leading nearly to the bottom of the evolution flask (B) and having itslower end bent upward and drawn out to small bore, so that the carbondioxide evolved from the limestone cannot bubble back into (b). Theevolution flask should preferably be a wide-mouthed Soxhlet extractionflask of about 150 cc. capacity because of the ease with which tubesand stoppers may be fitted into the neck of a flask of this type. Theflask should be fitted with a two-hole rubber stopper. The condenser(C) may consist of a tube with two or three large bulbs blown init, for use as an air-cooled condenser, or it may be a smallwater-jacketed condenser. The latter is to be preferred if a number ofdeterminations are to be made in succession.

A glass delivery tube (c) leads from the condenser to the small U-tube(D) containing some glass beads or small pieces of glass rod and 3 cc.of a saturated solution of silver sulphate, with 3 cc. of concentratedsulphuric acid (sp. gr. 1.84). The short rubber tubing (d) connectsthe first U-tube to a second U-tube (E) which is filled with smalldust-free lumps of dry calcium chloride, with a small, loose plug ofcotton at the top of each arm. Both tubes should be closed by corkstoppers, the tops of which are cut off level with, or preferablyforced a little below, the top of the U-tube, and then neatly sealedwith sealing wax.

The carbon dioxide may be absorbed in a tube containing soda lime(F) or in a Geissler bulb (F') containing a concentrated solutionof potassium hydroxide (Note 2). The tube (F) is a glass-stopperedside-arm U-tube in which the side toward the evolution flask and onehalf of the other side are filled with small, dust-free lumps of sodalime of good quality (Note 3). Since soda lime contains considerablemoisture, the other half of the right side of the tube is filled withsmall lumps of dry, dust-free calcium chloride to retain the moisturefrom the soda lime. Loose plugs of cotton are placed at the top ofeach arm and between the soda lime and the calcium chloride.

The Geissler bulb (F'), if used, should be filled with potassiumhydroxide solution (1 part of solid potassium hydroxide dissolved intwo parts of water) until each small bulb is about two thirds full(Note 4). A small tube containing calcium chloride is connected withthe Geissler bulb proper by a ground joint and should be wired to thebulb for safety. This is designed to retain any moisture from thehydroxide solution. A piece of clean, fine copper wire is so attachedto the bulb that it can be hung from the hook above a balance pan, orother support.

The small bottle (G) with concentrated sulphuric acid (sp. gr. 1.84)is so arranged that the tube (f) barely dips below the surface. Thiswill prevent the absorption of water vapor by (F) or (F') and servesas an aid in regulating the flow of air through the apparatus. (H) isan aspirator bottle of about four liters capacity, filled with water;(k) is a safety tube and a means of refilling (H); (h) is a screwclamp, and (K) a U-tube filled with soda lime.

[Note 1: The air current, which is subsequently drawn through theapparatus, to sweep all of the carbon dioxide into the absorptionapparatus, is likely to carry with it some hydrochloric acid fromthe evolution flask. This acid is retained by the silver sulphatesolution. The addition of concentrated sulphuric acid to this solutionreduces its vapor pressure so far that very little water is carried onby the air current, and this slight amount is absorbed by the calciumchloride in (E). As the calcium chloride frequently contains a smallamount of a basic material which would absorb carbon dioxide, it isnecessary to pass carbon dioxide through (E) for a short time and thendrive all the gas out with a dry air current for thirty minutes beforeuse.]

[Note 2: Soda-lime absorption tubes are to be preferred if asatisfactory quality of soda lime is available and the number ofdeterminations to be made successively is small. The potash bulbs willusually permit of a larger number of successive determinations withoutrefilling, but they require greater care in handling and in theanalytical procedure.]

[Note 3: Soda lime is a mixture of sodium and calcium hydroxides. Bothcombine with carbon dioxide to form carbonates, with the evolutionof water. Considerable heat is generated by the reaction, and thetemperature of the tube during absorption serves as a rough index ofthe progress of the reaction through the mass of soda lime.

It is essential that soda lime of good quality for analytical purposesshould be used. The tube should not contain dust, as this is likely tobe swept away.]

[Note 4: The solution of the hydroxide for use in the Geissler bulbmust be highly concentrated to insure complete absorption of thecarbon dioxide and also to reduce the vapor pressure of the solution,thus lessening the danger of loss of water with the air which passesthrough the bulbs. The small quantity of moisture which is thencarried out of the bulbs is held by the calcium chloride in theprolong tube. The best form of absorption bulb is that to which theprolong tube is attached by a ground glass joint.

After the potassium hydroxide is approximately half consumed in thefirst bulb of the absorption apparatus, potassium bicarbonate isformed, and as it is much less soluble than the carbonate, it oftenprecipitates. Its formation is a warning that the absorbing power ofthe hydroxide is much diminished.]

!The Analysis!

PROCEDURE.-- Weigh out into the flask (B) about 1 gram of limestone.Cover it with 15 cc. of water. Weigh the absorption apparatus (F)or (F') accurately after allowing it to stand for 30 minutes in thebalance case, and wiping it carefully with a lintless cloth, takingcare to handle it as little as possible after wiping (Note 1). Connectthe absorption apparatus with (e) and (f). If a soda-lime tube isused, be sure that the arm containing the soda lime is next the tube(E) and that the glass stopcocks are open.

To be sure that the whole apparatus is airtight, disconnect the rubbertube from the flask (A), making sure that the tubes (a) and (b) do notcontain any hydrochloric acid, close the pinchcocks (a) and (k) andopen (h). No bubbles should pass through (D) or (G) after a fewseconds. When assured that the fittings are tight, close (h) and open(a) cautiously to admit air to restore atmospheric pressure. Thisprecaution is essential, as a sudden inrush of air will project liquidfrom (D) or (F'). Reconnect the rubber tube with the flask (A).Open the pinchcocks (a) and (k) and blow over about 10 cc. of thehydrochloric acid from (A) into (B). When the action of the acidslackens, blow over (slowly) another 10 cc.

The rate of gas evolution should not exceed for more than a fewseconds that at which about two bubbles per second pass through (G)(Note 2). Repeat the addition of acid in small portions until theaction upon the limestone seems to be at an end, taking care to close(a) after each addition of acid (Note 3). Disconnect (A) and connectthe rubber tubing with the soda-lime tube (K) and open (a). Then close(k) and open (h), regulating the flow of water from (H) in such a waythat about two bubbles per second pass through (G). Place a smallflame under (B) and !slowly! raise the contents to boiling and boilfor three minutes. Then remove the burner from under (B) and continueto draw air through the apparatus for 20-30 minutes, or until (H)is emptied (Note 4). Remove the absorption apparatus, closing thestopcocks on (F) or stoppering the open ends of (F'), leave theapparatus in the balance case for at least thirty minutes, wipe itcarefully and weigh, after opening the stopcocks (or removing plugs).The increase in weight is due to absorption of CO_{2}, from which itspercentage in the sample may be calculated.

After cleaning (B) and refilling (H), the apparatus is ready for theduplicate analysis.

[Note 1: The absorption tubes or bulbs have large surfaces on whichmoisture may collect. By allowing them to remain in the balance casefor some time before weighing, the amount of moisture absorbed on thesurface is as nearly constant as practicable during two weighings, anda uniform temperature is also assured. The stopcocks of the U-tubeshould be opened, or the plugs used to close the openings of theGeissler bulb should be removed before weighing in order that the aircontents shall always be at atmospheric pressure.]

[Note 2: If the gas passes too rapidly into the absorption apparatus,some carbon dioxide may be carried through, not being completelyretained by the absorbents.]

[Note 3: The essential ionic changes involved in this procedure arethe following: It is assumed that the limestone, which is typified bycalcium carbonate, is very slightly soluble in water, and the ionsresulting are Ca^{++} and CO_{3}^{--}. In the presence of H^{+} ionsof the mineral acid, the CO_{3}^{--} ions form [H_{2}CO_{3}]. Thisis not only a weak acid which, by its formation, diminishes theconcentration of the CO_{3}^{--} ions, thus causing more of thecarbonate to dissolve to re-establish equilibrium, but it is also anunstable compound and breaks down into carbon dioxide and water.]

[Note 4: Carbon dioxide is dissolved by cold water, but the gas isexpelled by boiling, and, together with that which is distributedthrough the apparatus, is swept out into the absorption bulb by thecurrent of air. This air is purified by drawing it through the tube(K) containing soda lime, which removes any carbon dioxide which maybe in it.]

DETERMINATION OF LEAD, COPPER, IRON, AND ZINC IN BRASS

ELECTROLYTIC SEPARATIONS

!General Discussion!

When a direct current of electricity passes from one electrode toanother through solutions of electrolytes, the individual ions presentin these solutions tend to move toward the electrode of oppositeelectrical charge to that which each ion bears, and to be dischargedby that electrode. Whether or not such discharge actually occurs inthe case of any particular ion depends upon the potential (voltage) ofthe current which is passing through the solution, since for each ionthere is, under definite conditions, a minimum potential below whichthe discharge of the ion cannot be effected. By taking advantageof differences in discharge-potentials, it is possible to effectseparations of a number of the metallic ions by electrolysis, and atthe same time to deposit the metals in forms which admit of directweighing. In this way the slower procedures of precipitation andfiltration may frequently be avoided. The following paragraphs presenta brief statement of the fundamental principles and conditionsunderlying electro-analysis.

The total energy of an electric current as it passes through asolution is distributed among three factors, first, its potential,which is measured in volts, and corresponds to what is called "head"in a stream of water; second, current strength, which is measuredin amperes, and corresponds to the volume of water passing across-section of a stream in a given time interval; and third, theresistance of the conducting medium, which is measured in ohms. Therelation between these three factors is expressed by Ohm's law,namely, that !I = E/R!, when I is current strength, E potential, and Rresistance. It is plain that, for a constant resistance, thestrength of the current and its potential are mutually and directlyinterdependent.

As already stated, the applied electrical potential determines whetheror not deposition of a metal upon an electrode actually occurs. Thecurrent strength determines the rate of deposition and the physicalcharacteristics of the deposit. The resistance of the solution isgenerally so small as to fall out of practical consideration.

Approximate deposition-potentials have been determined for a numberof the metallic elements, and also for hydrogen and some of theacid-forming radicals. The values given below are those requiredfor deposition from normal solutions at ordinary temperatureswith reference to a hydrogen electrode. They must be regarded asapproximate, since several disturbing factors and some secondaryreactions render difficult their exact application under theconditions of analysis. They are:

From these data it is evident that in order to deposit copper from anormal solution of copper sulphate a minimum potential equal to thealgebraic sum of the deposition-potentials of copper ions and sulphateions must be applied, that is, +1.56 volts. The deposition of zincfrom a solution of zinc sulphate would require +2.67 volts, but, sincethe deposition of hydrogen from sulphuric acid solution requires only+1.90 volts, the quantitative deposition of zinc by electrolysis froma sulphuric acid solution of a zinc salt is not practicable. On theother hand, silver, if present in a solution of copper sulphate, woulddeposit with the copper.

The foregoing examples suffice to illustrate the application of theprinciple of deposition potentials, but it must further be notedthat the values stated apply to normal solutions of the compounds inquestion, that is, to solutions of considerable concentrations. As theconcentration of the ions diminishes, and hence fewer ions approachthe electrodes, somewhat higher voltages are required to attract anddischarge them. From this it follows that the concentrations should bekept as high as possible to effect complete deposition in the leastpracticable time, or else the potentials applied must be progressivelyincreased as deposition proceeds. In practice, the desired result isobtained by starting with small volumes of solution, using as large anelectrode surface as possible, and by stirring the solution to bringthe ions in contact with the electrodes. This is, in general, a moreconvenient procedure than that of increasing the potential of thecurrent during electrolysis, although that method is also used.

As already stated, those ions in a solution of electrolytes will firstbe discharged which have the lowest deposition potentials, and solong as these ions are present around the electrode in considerableconcentration they, almost alone, are discharged, but, as theirconcentration diminishes, other ions whose deposition potentials arehigher but still within that of the current applied, will also beginto separate. For example, from a nitric acid solution of coppernitrate, the copper ions will first be discharged at the cathode, butas they diminish in concentration hydrogen ions from the acid (orwater) will be also discharged. Since the hydrogen thus liberated is areducing agent, the nitric acid in the solution is slowly reduced toammonia, and it may happen that if the current is passed through for along time, such a solution will become alkaline. Oxygen is liberatedat the anode, but, since there is no oxidizable substance presentaround that electrode, it escapes as oxygen gas. It should be notedthat, in general, the changes occurring at the cathode are reductions,while those at the anode are oxidations.

For analytical purposes, solutions of nitrates or sulphates of themetals are preferable to those of the chlorides, since liberatedchlorine attacks the electrodes. In some cases, as for example, thatof silver, solution of salts forming complex ions, like that ofthe double cyanide of silver and potassium, yield better metallicdeposits.

Most metals are deposited as such upon the cathode; a few, notablylead and manganese, separate in the form of dioxides upon the anode.It is evidently important that the deposited material should be sofirmly adherent that it can be washed, dried, and weighed withoutloss in handling. To secure these conditions it is essential that thecurrent density (that is, the amount of current per unit of area ofthe electrodes) shall not be too high. In prescribing analyticalconditions it is customary to state the current strength in "normaldensities" expressed in amperes per 100 sq. cm. of electrode surface,as, for example, "N.D_{100} = 2 amps."

If deposition occurs too rapidly, the deposit is likely to be spongyor loosely adherent and falls off on subsequent treatment. This placesa practical limit to the current density to be employed, for a givenelectrode surface. The cause of the unsatisfactory character ofthe deposit is apparently sometimes to be found in the coincidentliberation of considerable hydrogen and sometimes in the failure ofthe rapidly deposited material to form a continuous adherent surface.The effect of rotating electrodes upon the character of the deposit isreferred to below.

The negative ions of an electrolyte are attracted to the anode and aredischarged on contact with it. Anions such as the chloride ion yieldchlorine atoms, from which gaseous chlorine molecules are formedand escape. The radicals which compose such ions as NO_{3}^{-} orSO_{4}^{--} are not capable of independent existence after discharge,and break down into oxygen and N_{2}O_{5} and SO_{3} respectively. Theoxygen escapes and the anhydrides, reacting with water, re-form nitricand sulphuric acids.

The law of Faraday expresses the relation between current strength andthe quantities of the decomposition products which, under constantconditions, appear at the electrodes, namely, that a given quantityof electricity, acting for a given time, causes the separation ofchemically equivalent quantities of the various elements or radicals.For example, since 107.94 grams of silver is equivalent to 1.008 gramsof hydrogen, and that in turn to 8 grams of oxygen, or 31.78 grams ofcopper, the quantity of electricity which will cause the deposit of107.94 grams of silver in a given time will also separate the weightsjust indicated of the other substances. Experiments show that acurrent of one ampere passing for one second, i.e., a coulomb ofelectricity, causes the deposition of 0.001118 gram of silver from anormal solution of a silver salt. The number of coulombs required todeposit 107.94 grams is 107.94/0.001118 or 96,550 and the same numberof coulombs will also cause the separation of 1.008 grams of hydrogen,8 grams of oxygen or 31.78 grams of copper. While it might at firstappear that Faraday's law could thus be used as a basis for thecalculation of the time required for the deposition of a givenquantity of an electrolyte from solution, it must be remembered thatthe law expresses what occurs when the concentration of the ions inthe solution is kept constant, as, for example, when the anode ina silver salt solution is a plate of metallic silver. Under theconditions of electro-analysis the concentration of the ions isconstantly diminishing as deposition proceeds and the time actuallyrequired for complete deposition of a given weight of material bya current of constant strength is, therefore, greater than thatcalculated on the basis of the law as stated above.

The electrodes employed in electro-analysis are almost exclusivelyof platinum, since that metal alone satisfactorily resists chemicalaction of the electrolytes, and can be dried and weighed withoutchange in composition. The platinum electrodes may be used in theform of dishes, foil or gauze. The last, on account of the ease ofcirculation of the electrolyte, its relatively large surface inproportion to its weight and the readiness with which it can be washedand dried, is generally preferred.

Many devices have been described by the use of which the electrodeupon which deposition occurs can be mechanically rotated. This has aneffect parallel to that of greatly increasing the electrode surfaceand also provides a most efficient means of stirring the solution.With such an apparatus the amperage may be increased to 5 or even 10amperes and depositions completed with great rapidity and accuracy. Itis desirable, whenever practicable, to provide a rotating or stirringdevice, since, for example, the time consumed in the deposition of theamount of copper usually found in analysis may be reduced from the20 to 24 hours required with stationary electrodes, and unstirredsolutions, to about one half hour.

DETERMINATION OF COPPER AND LEAD

PROCEDURE.--Weigh out two portions of about 0.5 gram each (Note 1)into tall, slender lipless beakers of about 100 cc. capacity. Dissolvethe metal in a solution of 5 cc. of dilute nitric acid (sp. gr. 1.20)and 5 cc. of water, heating gently, and keeping the beaker covered.When the sample has all dissolved (Note 2), wash down the sides of thebeaker and the bottom of the watch-glass with water and dilute thesolution to about 50 cc. Carefully heat to boiling and boil for aminute or two to expel nitrous fumes.

Meanwhile, four platinum electrodes, two anodes and two cathodes,should be cleaned by dipping in dilute nitric acid, washing with waterand finally with 95 per cent alcohol (Note 3). The alcohol may beignited and burned off. The electrodes are then cooled in a desiccatorand weighed. Connect the electrodes with the binding posts (or otherdevice for connection with the electric circuit) in such a way thatthe copper will be deposited upon the electrode with the largersurface, which is made the cathode. The beaker containing the solutionshould then be raised into place from below the electrodes until thelatter reach nearly to the bottom of the beaker. The support for thebeaker must be so arranged that it can be easily raised or lowered.

If the electrolytic apparatus is provided with a mechanism for therotation of the electrode or stirring of the electrolyte, proceed asfollows: Arrange the resistance in the circuit to provide a directcurrent of about one ampere. Pass this current through the solutionto be electrolyzed, and start the rotating mechanism. Keep the beakercovered as completely as possible, using a split watch-glass (or otherdevice) to avoid loss by spattering. When the solution is colorless,which is usually the case after about 35 minutes, rinse off the coverglass, wash down the sides of the beaker, add about 0.30 gram of ureaand continue the electrolysis for another five minutes (Notes 4 and5).

If stationary electrodes are employed, the current strength should beabout 0.1 ampere, which may, after 12 to 15 hours, be increased to 0.2ampere. The time required for complete deposition is usually from 20to 24 hours. It is advisable to add 5 cc. of nitric acid (sp. gr. 1.2)if the electrolysis extends over this length of time. No urea is addedin this case.

When the deposition of the copper appears to be complete, stop therotating mechanism and slowly lower the beaker with the left hand,directing at the same time a stream of water from a wash bottle onboth electrodes. Remove the beaker, shut off the current, and, ifnecessary, complete the washing of the electrodes (Note 6). Rinse theelectrodes cautiously with alcohol and heat them in a hot closet untilthe alcohol has just evaporated, but no longer, since the copper islikely to oxidize at the higher temperature. (The alcohol may beremoved by ignition if care is taken to keep the electrodes in motionin the air so that the copper deposit is not too strongly heated atany one point.)

Test the solution in the beaker for copper as follows, rememberingthat it is to be used for subsequent determinations of iron and zinc:Remove about 5 cc. and add a slight excess of ammonia. Compare themixture with some distilled water, holding both above a white surface.The solution should not show any tinge of blue. If the presence ofcopper is indicated, add the test portion to the main solution,evaporate the whole to a volume of about 100 cc., and againelectrolyze with clean electrodes (Note 7).

After cooling the electrodes in a desiccator, weigh them and from theweight of copper on the cathode and of lead dioxide (PbO_{2}) on theanode, calculate the percentage of copper (Cu) and of lead (Pb) in thebrass.

[Note 1: It is obvious that the brass taken for analysis should beuntarnished, which can be easily assured, when wire is used, byscouring with emery. If chips or borings are used, they should be wellmixed, and the sample for analysis taken from different parts of themixture.]

[Note 2: If a white residue remains upon treatment of the alloy withnitric acid, it indicates the presence of tin. The material is not,therefore, a true brass. This may be treated as follows: Evaporate thesolution to dryness, moisten the residue with 5 cc. of dilute nitricacid (sp. gr. 1.2) and add 50 cc. of hot water. Filter off themeta-stannic acid, wash, ignite in porcelain and weigh as SnO_{2}.This oxide is never wholly free from copper and must be purified foran exact determination. If it does not exceed 2 per cent of the alloy,the quantity of copper which it contains may usually be neglected.]

[Note 3: The electrodes should be freed from all greasy matter beforeusing, and those portions upon which the metal will deposit should notbe touched with the fingers after cleaning.]

[Note 4: Of the ions in solution, the H^{+}, Cu^{++}, Zn^{++}, andFe^{+++} ions tend to move toward the cathode. The NO_{3}^{-} ions andthe lead, probably in the form of PbO_{2}^{--} ions, move toward theanode. At the cathode the Cu^{++} ions are discharged and plate out asmetallic copper. This alone occurs while the solution is relativelyconcentrated. Later on, H^{+} ions are also discharged. In thepresence of considerable quantities of H^{+} ions, as in this acidsolution, no Zn^{++} or Fe^{+++} ions are discharged because of theirgreater deposition potentials. At the anode the lead is deposited asPbO_{2} and oxygen is evolved.

For the reasons stated on page 141 care must be taken that thesolution does not become alkaline if the electrolysis is longcontinued.]

[Note 5: Urea reacts with nitrous acid, which may be formed in thesolution as a result of the reducing action of the liberated hydrogen.Its removal promotes the complete precipitation of the copper. Thereaction is

CO(NH_{2})_{2} + 2HNO_{2} --> CO_{2} + 2N_{2} + 3H_{2}O.]

[Note 6: The electrodes must be washed nearly or quite free fromthe nitric acid solution before the circuit is broken to preventre-solution of the copper.

If several solutions are connected in the same circuit it is obviousthat some device must be used to close the circuit as soon as thebeaker is removed.]

[Note 7: The electrodes upon which the copper has been depositedmay be cleaned by immersion in warm nitric acid. To remove the leaddioxide, add a few crystals of oxalic acid to the nitric acid.]

DETERMINATION OF IRON

Most brasses contain small percentages of iron (usually not over 0.1per cent) which, unless removed, is precipitated as phosphate andweighed with the zinc.

PROCEDURE.--To the solution from the precipitation of copper andlead by electrolysis, add dilute ammonia (sp. gr. 0.96) until theprecipitate of zinc hydroxide which first forms re-dissolves, leavingonly a slight red precipitate of ferric hydroxide. Filter off theiron precipitate, using a washed filter, and wash five times with hotwater. Test a portion of the last washing with a dilute solution ofammonium sulphide to assure complete removal of the zinc.

The precipitate may then be ignited and weighed as ferric oxide, asdescribed on page 110.

Calculate the percentage of iron (Fe) in the brass.

DETERMINATION OF ZINC

PROCEDURE.--Acidify the filtrate from the iron determination withdilute nitric acid. Concentrate it to 150 cc. Add to the cold solutiondilute ammonia (sp. gr. 0.96) cautiously until it barely smells ofammonia; then add !one drop! of a dilute solution of litmus (Note 1),and drop in, with the aid of a dropper, dilute nitric acid until theblue of the litmus just changes to red. It is important that thispoint should not be overstepped. Heat the solution nearly to boilingand pour into it slowly a filtered solution of di-ammonium hydrogenphosphate[1] containing a weight of the phosphate about equalto twelve times that of the zinc to be precipitated. (For thiscalculation the approximate percentage of zinc is that found bysubtracting the sum of the percentages of the copper, lead and ironfrom 100 per cent.) Keep the solution just below boiling for fifteenminutes, stirring frequently (Note 2). If at the end of this time theamorphous precipitate has become crystalline, allow the solution tocool for about four hours, although a longer time does no harm (Note3), and filter upon an asbestos filter in a porcelain Gooch crucible.The filter is prepared as described on page 103, and should be driedto constant weight at 105 deg.C.

[Footnote 1: The ammonium phosphate which is commonly obtainablecontains some mono-ammonium salt, and this is not satisfactory as aprecipitant. It is advisable, therefore, to weigh out the amount ofthe salt required, dissolve it in a small volume of water, add a dropof phenolphthalein solution, and finally add dilute ammonium hydroxidesolution cautiously until the solution just becomes pink, but do notadd an excess.]

Wash the precipitate until free from sulphates with a warm 1 per centsolution of the di-ammonium phosphate, and then five times with 50 percent alcohol (Note 4). Dry the crucible and precipitate for an hour at105 deg.C., and finally to constant weight (Note 5). The filtrate shouldbe made alkaline with ammonia and tested for zinc with a few drops ofammonium sulphide, allowing it to stand (Notes 6, 7 and 8).

From the weight of the zinc ammonium phosphate (ZnNH_{4}PO_{4})calculate the percentage of the zinc (Zn) in the brass.

[Note 1: The zinc ammonium phosphate is soluble both in acids and inammonia. It is, therefore, necessary to precipitate the zinc in anearly neutral solution, which is more accurately obtained by addinga drop of a litmus solution to the liquid than by the use of litmuspaper.]

[Note 2: The precipitate which first forms is amorphous, and may havea variable composition. On standing it becomes crystalline and thenhas the composition ZnNH_{4}PO_{4}. The precipitate then settlesrapidly and is apt to occasion "bumping" if the solution is heated toboiling. Stirring promotes the crystallization.]

[Note 3: In a carefully neutralized solution containing a considerableexcess of the precipitant, and also ammonium salts, the separationof the zinc is complete after standing four hours. The ionic changesconnected with the precipitation of the zinc as zinc ammoniumphosphate are similar to those described for magnesium ammoniumphosphate, except that the zinc precipitate is soluble in an excess ofammonium hydroxide, probably as a result of the formation of complexions of the general character Zn(NH_{3})_{4}^{++}.]

[Note 4: The precipitate is washed first with a dilute solution of thephosphate to prevent a slight decomposition of the precipitate (as aresult of hydrolysis) if hot water alone is used. The alcohol is addedto the final wash-water to promote the subsequent drying.]

[Note 5: The precipitate may be ignited and weighed asZn_{2}P_{2}O_{7}, by cautiously heating the porcelain Gooch cruciblewithin a nickel or iron crucible, used as a radiator. The heatingmust be very slow at first, as the escaping ammonia may reduce theprecipitate if it is heated too quickly.]

[Note 6: If the ammonium sulphide produced a distinct precipitate,this should be collected on a small filter, dissolved in a few cubiccentimeters of dilute nitric acid, and the zinc reprecipitated asphosphate, filtered off, dried, and weighed, and the weight added tothat of the main precipitate.]

[Note 7: It has been found that some samples of asbestos are actedupon by the phosphate solution and lose weight. An error from thissource may be avoided by determining the weight of the crucibleand filter after weighing the precipitate. For this purpose theprecipitate may be dissolved in dilute nitric acid, the asbestoswashed thoroughly, and the crucible reweighed.]

[Note 8. The details of this method of precipitation of zinc are fullydiscussed in an article by Dakin, !Ztschr. Anal. Chem.!, 39 (1900),273.]

DETERMINATION OF SILICA IN SILICATES

Of the natural silicates, or artificial silicates such as slags andsome of the cements, a comparatively few can be completely decomposedby treatment with acids, but by far the larger number require fusionwith an alkaline flux to effect decomposition and solutionfor analysis. The procedure given below applies to silicatesundecomposable by acids, of which the mineral feldspar is taken as atypical example. Modifications of the procedure, which are applicableto silicates which are completely or partially decomposable by acids,are given in the Notes on page 155.

PREPARATION OF THE SAMPLE

Grind about 3 grams of the mineral in an agate mortar (Note 1) untilno grittiness is to be detected, or, better, until it will entirelypass through a sieve made of fine silk bolting cloth. The sieve may bemade by placing a piece of the bolting cloth over the top of a smallbeaker in which the ground mineral is placed, holding the cloth inplace by means of a rubber band below the lip of the beaker. Byinverting the beaker over clean paper and gently tapping it, the fineparticles pass through the sieve, leaving the coarser particles withinthe beaker. These must be returned to the mortar and ground, and theprocess of sifting and grinding repeated until the entire samplepasses through the sieve.

[Note 1: If the sample of feldspar for analysis is in the massive orcrystalline form, it should be crushed in an iron mortar until thepieces are about half the size of a pea, and then transferred to asteel mortar, in which they are reduced to a coarse powder. A woodenmallet should always be used to strike the pestle of the steel mortar,and the blows should not be sharp.

It is plain that final grinding in an agate mortar must be continueduntil the whole of the portion of the mineral originally taken hasbeen ground so that it will pass the bolting cloth, otherwise thesifted portion does not represent an average sample, the softeringredients, if foreign matter is present, being first reduced topowder. For this reason it is best to start with not more than thequantity of the feldspar needed for analysis. The mineral must bethoroughly mixed after the grinding.]

FUSION AND SOLUTION

PROCEDURE.--Weigh into platinum crucibles two portions of the groundfeldspar of about 0.8 gram each. Weigh on rough balances two portionsof anhydrous sodium carbonate, each amounting to about six times theweight of the feldspar taken for analysis (Note 1). Pour about threefourths of the sodium carbonate into the crucible, place the latter ona piece of clean, glazed paper, and thoroughly mix the substance andthe flux by carefully stirring for several minutes with a dry glassrod, the end of which has been recently heated and rounded in a flameand slowly cooled. The rod may be wiped off with a small fragment offilter paper, which may be placed in the crucible. Place the remainingfourth of the carbonate on the top of the mixture. Cover the crucible,heat it to dull redness for five minutes, and then gradually increasethe heat to the full capacity of a Bunsen or Tirrill burner fortwenty minutes, or until a quiet, liquid fusion is obtained (Note 2).Finally, heat the sides and cover strongly until any material whichmay have collected upon them is also brought to fusion.

Allow the crucible to cool, and remove the fused mass as directed onpage 116. Disintegrate the mass by placing it in a previously preparedmixture of 100 cc. of water and 50 cc. of dilute hydrochloric acid(sp. gr. 1.12) in a covered casserole (Note 3). Clean the crucible andlid by means of a little hydrochloric acid, adding this acid to themain solution (Notes 4 and 5).

[Note 1: Quartz, and minerals containing very high percentages ofsilica, may require eight or ten parts by weight of the flux to insurea satisfactory decomposition.]

[Note 2: During the fusion the feldspar, which, when pure, is asilicate of aluminium and either sodium or potassium, but usuallycontains some iron, calcium, and magnesium, is decomposed by thealkaline flux. The sodium of the latter combines with the silicic acidof the silicate, with the evolution of carbon dioxide, while about twothirds of the aluminium forms sodium aluminate and the remainderis converted into basic carbonate, or the oxide. The calcium andmagnesium, if present, are changed to carbonates or oxides.

The heat is applied gently to prevent a too violent reaction whenfusion first takes place.]

[Note 3: The solution of a silicate by a strong acid is the result ofthe combination of the H^{+} ions of the acid and the silicate ionsof the silicate to form a slightly ionized silicic acid. As aconsequence, the concentration of the silicate ions in the solution isreduced nearly to zero, and more silicate dissolves to re-establishthe disturbed equilibrium. This process repeats itself until all ofthe silicate is brought into solution.

Whether the resulting solution of the silicate contains ortho-silicicacid (H_{4}SiO_{4}) or whether it is a colloidal solution of someother less hydrated acid, such as meta-silicic acid (H_{2}SiO_{3}),is a matter that is still debatable. It is certain, however, that thegelatinous material which readily separates from such solutions is ofthe nature of a hydrogel, that is, a colloid which is insoluble inwater. This substance when heated to 100 deg.C., or higher, is completelydehydrated, leaving only the anhydride, SiO_{2}. The changes may berepresented by the equation:

SiO_{3}^{--} + 2H^{+} --> [H_{2}SiO_{3}] --> H_{2}O + SiO_{2}.]

[Note 4: A portion of the fused mass is usually projected upward bythe escaping carbon dioxide during the fusion. The crucible musttherefore be kept covered as much as possible and the lid carefullycleaned.]

[Note 5: A gritty residue remaining after the disintegration ofthe fused mass by acid indicates that the substance has been butimperfectly decomposed. Such a residue should be filtered, washed,dried, ignited, and again fused with the alkaline flux; or, if thequantity of material at hand will permit, it is better to reject theanalysis, and to use increased care in grinding the mineral and inmixing it with the flux.]

DEHYDRATION AND FILTRATION

PROCEDURE.--Evaporate the solution of the fusion to dryness, stirringfrequently until the residue is a dry powder. Moisten the residue with5 cc. of strong hydrochloric acid (sp. gr. 1.20) and evaporate againto dryness. Heat the residue for at least one hour at a temperatureof 110 deg.C. (Note 1). Again moisten the residue with concentratedhydrochloric acid, warm gently, making sure that the acid comes intocontact with the whole of the residue, dilute to about 200 cc. andbring to boiling. Filter off the silica without much delay (Note 2),and wash five times with warm dilute hydrochloric acid (one partdilute acid (1.12 sp. gr.) to three parts of water). Allow the filterto drain for a few moments, then place a clean beaker below the funneland wash with water until free from chlorides, discarding thesewashings. Evaporate the original filtrate to dryness, dehydrate at110 deg.C. for one hour (Note 3), and proceed as before, using a secondfilter to collect the silica after the second dehydration. Wash thisfilter with warm, dilute hydrochloric acid (Note 4), and finally withhot water until free from chlorides.

[Note 1: The silicic acid must be freed from its combination witha base (sodium, in this instance) before it can be dehydrated.The excess of hydrochloric acid accomplishes this liberation. Bydisintegrating the fused mass with a considerable volume of diluteacid the silicic acid is at first held in solution to a large extent.Immediate treatment of the fused mass with strong acid is likelyto cause a semi-gelatinous silicic acid to separate at once and toinclose alkali salts or alumina.

A flocculent residue will often remain after the decomposition of thefused mass is effected. This is usually partially dehydrated silicicacid and does not require further treatment at this point. Theprogress of the dehydration is indicated by the behavior of thesolution, which as evaporation proceeds usually gelatinizes. On thisaccount it is necessary to allow the solution to evaporate on a steambath, or to stir it vigorously, to avoid loss by spattering.]

[Note 2: To obtain an approximately pure silica, the residue afterevaporation must be thoroughly extracted by warming with hydrochloricacid, and the solution freely diluted to prevent, as far as possible,the inclosure of the residual salts in the particles of silica. Thefiltration should take place without delay, as the dehydrated silicaslowly dissolves in hydrochloric acid on standing.]

[Note 3: It has been shown by Hillebrand that silicic acid cannot becompletely dehydrated by a single evaporation and heating, nor byseveral such treatments, unless an intermediate filtration of thesilica occurs. If, however, the silica is removed and the filtratesare again evaporated and the residue heated, the amount of silicaremaining in solution is usually negligible, although severalevaporations and filtrations are required with some silicates toinsure absolute accuracy.

It is probable that temperatures above 100 deg.C. are not absolutelynecessary to dehydrate the silica; but it is recommended, as tendingto leave the silica in a better condition for filtration than whenthe lower temperature of the water bath is used. This, and many otherpoints in the analysis of silicates, are fully discussed by Dr.Hillebrand in the admirable monograph on "The Analysis of Silicate andCarbonate Rocks," Bulletin No. 700 of the United States GeologicalSurvey.

The double evaporation and filtration spoken of above are essentialbecause of the relatively large amount of alkali salts (sodiumchloride) present after evaporation. For the highest accuracy in thedetermination of silica, or of iron and alumina, it is also necessaryto examine for silica the precipitate produced in the filtrate byammonium hydroxide by fusing it with acid potassium sulphate and